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Models of Acids and Bases - Coggle Diagram
Models of Acids and Bases
Properties of acids and bases
Acids
mostly colourless except organic acids
sour taste
soluble in water
pH range 1-7
turn blue litmus red
feel like water
Bases
colourless except hydroxides of iron and copper
bitter taste
feel slippery
some bases soluble in water
pH range >7
turn red litmus blue
acid + base --> salt + water
neutralisation reaction
The Arrhenius theory
attempted to explain acids and bases in terms of the particles they produced in aqueous solution
PROPOSED
An acid was a substance that ionised in solution to produce hydrogen ions
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A base was a substance that in solution produced hydroxide ions
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If HA is used to represent an acid
HA (aq) --> H+ (aq) + A-
XOH is used to represent a base
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XOH (aq) --> X+ (aq) + OH- (aq)
The theory was consistent with all acids and bases except could not explain why ammonia is a base
The dilemma was that either ammonia was not a base or the definition was inaccurate or incomplete
This showed that the
original definition of an acid and a base needed to be redefined
Limitations
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The nature and role of the solvent was not considered
An acid should exhibit acidic properties in any solvent
According to the model, all salts produced by reactions of an acid and a base should be neutral
This is not the case in reality
Not all bases are soluble
Thus OH- ions cannot be produced
when bases are dissolved in non-aqueous solvents, no hydroxide ions are present in the resulting solution
The Bronsted-Lowry theory
redefined an acid and a base to take ammonia into account
PROPOSED
A acid is a substance that
donates one or more hydrogen ions
A base is a substance that
accepts one or more protons
Bases become the mirror companion of acids
The definitions are broader than the Arrhenius definitions
Example
NH3 (aq) + H2O (l) --><-- NH4+ (aq) + OH- (aq)
Limitations
for a substance to be identified as an aid or base, a H+ must be transferred in the reaction
this model requires the presence of a solvent, which has a hydrogen attached to an oxygen or nitrogen
It does not explain the acid-base behaviour of substances in non-aqueous solvents where a proton is not involved
cannot explain reactions between acidic oxides, such as CO2, SO2, SO3
and basic oxides like CaO, BaO, MgO
allows the reactions of acids and bases to be reversible
HA + B --><-- HB+ + A-
HA is an acid as it is donating a proton to the base
in the reverse reaction, HB+ is the acid, as it is donating the proton to the A- base
The acid base pairs are called conjugate pairs
Example: HCl (aq) + H2O (l) --><-- H3O+ (aq) + Cl- (aq)
This example here shows the acid HCl donating a proton to form its conjugate base, Cl-. Water is accepting a proton to act as the base and form its conjugate acid, H3O+.
Lewis definition
PROPOSED
An acid is an electron pair acceptor
A base is an electron pair donor
Even broader definition than the Bronsted-Lowry theory
Does not require either a proton or a solvent
A Lewis acid is any atom, ion or molecule that can accept electrons
A Lewis base is any atom, ion, or molecule capable of donating electrons
Example
BF3 reacts with ammonia
pH
pH scale
shows degree of acidity and alkalinity of solutions
based of the concentration of hydrogen ions in solution
reflects importance of the Bronsted-Lowry definition of an acid as a proton donor
goes from 0-14
lower the pH, more acidic
high pH, more basic
at 25 degrees celsius, a substance of pH 7 is neutral
Measuring pH
pH probes
pH meters
universal indicator
produces several colour changes depending on the pH of solution
mixture of several indicators
several other indicators to test acids and bases
pH = -log10[H+]